Chapter 12

1. Chapter 12 Liquids and Solids

2. Sect. 12-1: Liquids • Properties of liquids • Definite volume • Takes shape of its container • Properties can be explained by the kinetic molecular theory (KMT) just like gases were explained • Fluid – has the ability to flow

3. Properties of liquids (con’t) • Relatively high density • Relative incompressibility • Ability to diffuse • Surface tension – force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size

4. More properties of liquids… • Capillary action – the attraction of the surface of a liquid to the surface of a solid • Vaporization – liquid changes to a gas • Evaporation – vaporization occurring at the surface of a non-boiling liquid • Freezing – liquid changing to a solid

5. Sect. 12-2: Solids • Properties of solids • Definite volume • Definite shape • Properties can be explained by the kinetic molecular theory (KMT) just like gases and liquids were explained

6. Two types of solids • Crystalline solids – made up of crystals • Crystals – particles are arranged in an orderly, geometric fashion • Amorphous solids – particles are arranged randomly

7. More properties of solids… • Definite melting point – temperature at which a crystalline solid turns into a liquid • Supercooled liquids – substances that retain certain liquid properties even at temperatures at which they appear to be solid (amorphous solids) • High density • Incompressibility • Very low rate of diffusion

8. Crystalline solids • Crystal structure – total 3D arrangement of particles of a crystal • Unit cell – smallest part of a crystal structure that shows the 3D pattern of the entire structure

9. Binding forces in crystals • Ionic crystals • Covalent network crystals • Metallic crystals • Covalent molecular crystals

10. Sect.12-3: Changes of State • Equilibrium – a dynamic condition in which two opposing changes occur at equal rates in a closed system • Phase – any part of a system that has uniform composition and properties

11. Equilibrium and changes of state • If you start out with a closed container of liquid, eventually some of it will evaporate. • Over time, some of the gas will condense. • If left for a long enough period of time, equilibrium will be reached for amount of substance condensing vs. evaporating. • Represented by the following equation: liquid + heat energy  vapor

12. System will stay at equilibrium until some outside force acts upon it. • Le Châtelier’s Principle – when a system at equilibrium is disturbed by an application of stress, it attains a new equilibrium position that minimizes the stress • Stress is typically a change in concentration, pressure, or temperature.

13. Equilibrium and temperature • What will happen to earlier discussed liquid/vapor system if temperature is increased? Which reaction will be favored? • What if temp is decreased? Which reaction is favored?

14. Equilibrium and concentration • What would happen to system if volume of container was suddenly increased? • Amount of vapor would now be in lower concentration and thus less condensation would be occurring and more evaporation occurring until system reaches equilibrium again. • At the new equilibrium, there fewer liquid molecules, but the same concentration of the gas due to larger container size. (equilibrium has shifted right)

15. Equilibrium and Vapor Pressure of Liquid • Equilibrium vapor pressure – pressure exerted by a vapor in equilibrium over a liquid at a given temperature (increases as temp increases) • Explained by KMT • Pg. 376 for illustration

16. Volatile vs. nonvolatile liquids • Volatile liquid – liquids which evaporate very easily due to weak forces of attraction between particles • Have a relatively high equilibrium vapor pressure • Nonvolatile liquids have stronger forces of attraction and lower equilibrium vapor pressures

17. Boiling • Boiling – conversion of a liquid to a gas at the surface and throughout the liquid • Boiling point – temperature at which the equilibrium vapor pressure is equal to the atmospheric pressure • Pressure affects the boiling point • Normal boiling point is at 1 atm pressure

18. During a phase change, such as vaporization, the temperature remains constant because any heat being added is going to changing the particles from liquid to gas • Molar heat of vaporization – amount of heat required to vaporize one mole of liquid at its boiling point • The higher it is, the stronger the bonds are that hold that liquid together

19. Freezing and Melting • Freezing point – temperature at which the solid and liquid are in equilibrium at 1 atm • Molar heat of fusion – amount of heat required to melt one mole of a solid at its melting point

20. Sublimation – change of state from a solid directly to a gas • Deposition – change of state from a gas directly to a solid • Both occur at low pressures and temperatures where a liquid cannot exist

21. Phase Diagrams • Phase diagram – a graph of pressure vs. temperature that shows the conditions under which the phases of a substance exist

23. Triple point – indicates the temperature and pressure at which the solid, liquid, and vapor can coexist at equilibrium • Critical point – indicates critical temp. and pressure • Critical temperature (tc)– temp above which a substance cannot exist as a liquid • Critical pressure (Pc) – lowest pressure at which the substance can exist as a liquid at the critical temperature

24. Sect. 12-4: Water • Water molecules have a bent structure • Solid or liquid water molecules are linked by hydrogen bonding • Ice is made up of water arranged in hexagon shapes; this allows for low density

25. Physical Properties of Water • Liquid water is odorless, tasteless, & transparent • Relatively high molar heat of fusion (6.009 kJ/mol) • Relatively high boiling point and molar heat of vaporization (40.79 kJ/mol) compared to other substances that have similar molar masses (because of hydrogen bonding)